The simplest chemical reactions are those that occur in
the gas phase in a single step, such as the transfer of a
chlorine atom from ClNO2 to NO to form NO2
This reaction can be understood by writing
the Lewis structures for all four components of the reaction.
Both NO and NO2 contain an odd number of
electrons. Both NO and NO2 can therefore combine
with a neutral chlorine atom to form a molecule in which all
of the electrons are paired. This reaction therefore involves
the transfer of a chlorine atom from one molecule to another,
as shown in the figure below.
The figure below combines a plot of the
disappearance of the ClNO2 consumed in this
reaction with a plot of the appearance of NO2
formed in the reaction.
One of the goals of collecting these data
is to describe the rate of reaction, which
is the rate at which the reactants are transformed into the
products of the reaction.
The mathematical equation that describes the rate of a
chemical reaction is called the rate law for
the reaction. The data in the figure above are consistent
with the following rate law for this reaction.
Rate = k(ClNO2)(NO)
According to this rate law, the rate at which ClNO2
and NO are converted into NO2 and ClNO is
proportional to the product of the concentrations of the two
reactants. Initially, the rate of reaction is fast. As the
reactants are converted into products, however, the ClNO2
and NO concentrations become smaller, and the reaction slows
We might expect the reaction to stop when it runs out of
either ClNO2 or NO. In practice, the reaction
stops before this happens.
This is a very fast reaction the concentration of ClNO2
drops by a factor of two in less than a second. And yet, no
matter how long we wait, some residual ClNO2 and
NO remains in the reaction flask.
The figure below divides the plot of the change in the
concentrations of NO2 and ClNO into a kinetic
region and an equilibrium region.
By definition, the kinetic region is the period during which
the concentrations of the components of the reaction are
constantly changing. The equilibrium region is the period
after which the reaction seems to stop, when there is no
further change in the concentrations of the components of the
Theory Model for Gas-Phase Reactions
The fact that the following reaction
seems to stop before all of the reactants
are consumed can be explained with a model of chemical
reactions known as the collision theory.
This model assumes that ClNO2 and NO molecules
must collide before a chlorine atom can be transferred from
one molecule to the other.
This assumption explains why the rate of the reaction is
proportional to the concentration of both ClNO2
Rate = k(ClNO2)(NO)
The number of collisions per second between ClNO2
and NO molecules depends on their concentrations. As ClNO2
and NO are consumed in the reaction, the number of collisions
per second between these molecules becomes smaller, and the
reaction slows down.
Suppose that we start with a mixture of ClNO2
and NO, but no NO2 or ClNO. The only reaction that
can occur at first is the transfer of a chlorine atom from
ClNO2 to NO.
Eventually, NO2 and ClNO build
up in the reaction flask, and these molecules begin to
collide as well. Collisions between these molecules can
result in the transfer of a chlorine atom in the opposite
The collision theory model of chemical
reactions assumes that the rate of a simple, one-step
reaction is proportional to the product of the concentrations
of the ions or molecules consumed in that reaction. The rate
of the forward reaction is therefore proportional to the
product of the concentrations of the two
Rateforward = kf(ClNO2)(NO)
The rate of the reverse reaction, on the other hand, is
proportional to the concentrations of the
"products" of the reaction.
Ratereverse = kr(NO2)(ClNO)
Initially, the rate of the forward reaction is much larger
than the rate of the reverse reaction, because the system
contains ClNO2 and NO, but virtually no NO2
As ClNO2 and NO are consumed,
the rate of the forward reaction slows down. At the same
time, NO2 and ClNO accumulate, and the reverse
reaction speeds up.
If the forward reaction gradually slows down and the
reverse reaction speeds up, the system eventually has to
reach a point at which the rates of the forward and reverse
reactions are the same.
At this point, the reaction will seem to
stop. ClNO2 and NO will be consumed in the forward
reaction at the rate at which they are produced in the
reverse reaction. The same thing will happen to NO2
and ClNO. When the rates of the forward and reverse reactions
are the same, there is no longer any change in the
concentrations of the reactants or products of the reaction.
In other words, the reaction is at equilibrium.
We can now see that there are two definitions of
1. A system in which there is no apparent change in the
concentrations of the reactants and products of a reaction.
2. A system in which the rates of the forward and reverse
reactions are equal.
The first definition is based on the results of
experiments that tell us that some reactions seem to stop
prematurely they reach a point at which no more reactants are
converted into products before the limiting reagent is
consumed. The other definition is based on a theoretical
model of chemical reactions that explains why reactions reach
We can now distinguish between reactions that go to
completion and those that reach equilibrium. Reactions that
aren't reversible, or that strongly favor the products, are
assumed to go to completion and are represented by equations
that contain a single arrow.
Reversible reactions that reach equilibrium
are indicated by a pair of arrows between the two sides of